The Mole and Avogadros Number

There are several reasons why this number is useful in chemistry. First, it is a large number that gives a useftd unit for counting particles as small as atoms. In 

The mass of 6.022 x 10 atoms of any element is the molar mass of that element, and its value in grams per mole is given in most periodic tables. The molar mass of silicon, for example, is greater than that of carbon because an atom of silicon is heavier than an atom of carbon. And the molar mass of each element takes into account the natural isotopic abundances. So the molar mass of carbon is 12.011 g/mol, reflecting the weighted average between the masses of and in the same way as we discussed in Chapter 2. 

The mole is the key to the macroscopic interpretation of a chemical equation. Consider the same equation we have been discussing  

If we want to read this in terms of moles, we can say two moles of H2 and one mole of O2 react to form two moles of H2O. Because each mole contains the same number of molecules, the 2 1 mole ratio between the reactants is the same as the 2 1 ratio for the numbers of molecules. Chemical equations and their stoichiometric coefficients always provide ratios of numbers of particles, not masses. 

Even the smallest samples we deal with in the laboratory contain enormous numbers of atoms, ions, or molecules. For example, a teaspoon of water (about 5 mL) contains 2 X 10 water molecules, a number so large it almost defies comprehension. Chemists therefore have devised a counting unit for describing such large numbers of atoms or molecules. 

In everyday life we use such familiar counting units as dozen (12 objects) and gross (144 objects). In chemistry the counting unit for numbers of atoms, ions, or molecules in a laboratory-size sample is the mole, abbreviated mol. One mole is the amount of matter that contains as many objects (atoms, molecules, or diatever other objects we are considering) as the number of atoms in exactly 12 g of isotopicaUy pure C. From experiments, scientists have determined this number to be 6.0221421 X 10, diich we will usually round to 6.02 X 10. Scientists call this value Avogadro s number, N, in honor 

Practice is the key to success in solving problems. As you practice, you can improve your skills by following these steps  

Step 1 Analyze the problem. Read the problem carefully. What does it say Draw a picture or diagram that will help you to visuahze the problem. Write down both the data you are given and the quantity you need to obtain (the unknown). 

Step 2 Develop a plan for solving the problem. Consider a possible path between the given information and the unknown. What principles or equations relate the known data to the unknown  

Chemists must sometimes calculate the percentage composition of a compound—that is, the percentage by mass contributed by each element in the substance. Forensic chemists, for example, can measure the percentage composition of an unknown powder and compare it with the percentage compositions of suspected substances (for example, sugar, salt, or cocaine) to identify the powder. 

Calculating the percentage composition of any element in a substance (sometimes called the elemental composition of a substance) is straightforward if the chemical formula is known. The calculation depends on the formula weight of the substance, the atomic weight of the element of interest, and the number of atoms of that element in the chemical formula  

Calculate the percentage of carbon, hydrogen, and oxygen (by mass) in C12H22O11. 

Let s examine this question using the problem-solving steps in the accompanying Strategies in Chemistry Problem Solving essay. 

Analyze We are given a chemical formula and asked to calculate the percentage by mass of each element. 

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